NOTES
1056
tration, ko, fails to obey the Arrhenius rate law most spectacularly of all, apparently rising to a maximum around 120” and falling to zero at 147”. This strange behavior is not due to changes in the solubility of oxygen because the solubility of oxygen m water is nearly temperature independent in the pertinent range of temperatures.6 More careful experimental work may show that k~ does obey the Arrhenius (and the Eyring) rate laws. The four points shown appear to represent scatter about a, line. Considering that each k~ is the slope of a line determined from only two experiments, that scatter may not be in excess of the uncertainty in the points. At all temperatures, k~ is considerably larger than the second-order rate constant for acid cleavage obtained in the absence of oxygen. At 11l0, for example, k~ is 2.69 X 1. mole-’ sec. -I, while the second-order rate constant for acid cleavage in the absence of oxygen is 1.3 X 10+ 1. mole-’ set.-'. The reaction of t-butylmercuric iodide was not thoroughly investigated because of the solubility difficulties encountered, but there, also, oxygen had a very profound effect on the observed rates. For example, with 0.136 M perchloric acid and ampoules sealed in air, a first-order rate constant of 3.32 X 10” 1. mole-’ set.-' was obtained a t 100”. Under identical conditions but with each ampoule bubbled with oxygen-free nitrogen for one minute before sealing, the rate constant was 1.39 X lo-* 1. mole-’ sec.-1. The latter may still be higher than the oxygen-free rate because more stringent precautions had to be taken to reach a minimum rate a t higher temperature^.^
Discussion To our knowledge, there is no previous report of a reaction be tween organomercurials and oxygen, although it is well known that other organometallics react with oxygen and the reactions have been extensively studied.’ Equations 3-6 show an attractivc path by which the present reaction might proceed. RIIgI
several
+-+ -+
+
steps
ROOHgI
+ +
RO0:EIgI H e -+ ROOH IIgIe ROOH +decomposition products RHgI H g I e -+ RHg+ HgIz
+
(3)
(4) (5) (6)
The temperature dependence of the rate, particularly at, the lower acid concentrations, strongly suggests that the key steps, represented in eq. 3, are decidedly unusual in character. More mechanistic speculation does not seem in order a t this stage. Experimental The general ~xoceduresused in preparing solvents and reagents have been described previously,‘ as has the preparation of the si1bstrates.a All of the “aqueous” solutions referred to above actually contained 2% of methanol since the substrate8 were handled as stock solutions in that solvent. (6) H. A. Pray, C. Schweickert and B. Minnich, Ind. Eng. Chem.,
44, 1146 (1952). (7) H. Hock, 11. IZropf and F. Ernst, Angew. Chem., 11, 541 (1959).
Vol. 65
SOME THERMODYNAMIC PROPERTIES OF T H E SYSTEM PLUTONIUM CHLORIDESODIUM CHLORIDE FROM ELECTROMOTIVE FORCE DATA BY R. BENZAND J. A. LEARY University of CaliJornia, Los Alamos Scientific Laboratory, Lo8 Alamos, New Meiico Received November 18, 1960
Recently a study’ of the free energy of formation of PuC13in the binary system PuCl3-KC1 was made using potentiometric methods. This paper is a report on e.m.f. data as a function of temperature and composition obtained from galvanic cells of the type PU(ii,)/PuCla-NaClc I iq) /Cle(g)
( 1)
The cell reaction is Pu(liq)
+ 35 G ~ w P~C13(1iq, =
XI)
(2)
The thermodynamic formation quantities are derived and compared with those taken from the previous study. Experimental Materials.-PuC13 was prepared as previously described.2 The plutonium used for the metal electrodes was 99.8 i 0.1% pure by chemical analysis. Sodium chloride (Baker and Adamson, A.C.S. grade) was dried, melted in a quartz container under an atmosphere of hydrogen chloride, cast into a stick form and stored in a vacuum desiccator. The melting point of the sodium chloride was determined to be 801 f 2’ in a reement with the reported value.* Hydrogen chloride ?Matheson Co.) was dried with magnesium perchlorate. Argon (Linde Air Products) was dried with phosphorus pentoxide. Chlorine gas (Mathrson Go., 99.3y0 pure) was used without further treatment. The pressure of the chlorine gas, which was measured with an accuracy of i l mm. in the range 588 to 508 mm., was corrected to one atmosphere by assuming thc gas t o be ideal and adding ( - R T In p ) j 2 to the observed e.m .f . values. The chemical composition of each solution (25 to 75 g. total weight) was based upon the composition by weight of the component salts. Apparatus and Procedure.-The apparatus and procedure has been described.’ A single measurement usually extended over a period of 5 minutes. During this period the e.m.f. readings were verified to be reproducible after applying for two seconds an electrodepositing current of 15 ma. a t 2.5 v. and, again, after applying for two seconds an electrodissolving current. After a measurement was completed, adjustments were made to establish a new temperature equilibrium, shifted by 10 to 30°, increasing and decreasing at alternate measurements such that alternate equilibrium temperatures were approached from above and below. Such a series of otentiometric measurements was carried out on each of the aquid solutions4 in the temperature range 640 to 755” for the PuCla mole fractions 0.750, 0.700, 0.600 and 0.500. The cells remained reversible for periods of from 2 to 11 hours.
Results Calculations.-The notation employed by Wagner5 for the thermodynamic mixing properties is followed. The subscript 1 denotes liquid -PuCl3 and 2 denotes liquid sodium chloride. X1 and X Z (1) R. Bens, J . Phys. Chem., 66, 81 (1961). (2) R. Bens, Milton Kahn and J. A. Leary, ibid., 68, 1983 (1959). (3) National Bureau of Standards Circular No. 500, 1952, p. 804. (4) C. W. Bjorklund, J. G. Reavis, J. A. Leary and K. A. Walsh, J . Phys. Chem., 6 3 , 1774 (1959). (5) C. Wagner, “Thermodynamics of Alloys.” Addison-Wesley Press. Inc., New York, N. Y.,1952,Chap. 1.
NOTES
June, 1961 denote the mole fractions of PuCL and NaC1, respectively. The reference state for the relative partial quantities is that of the supercooled liquid PuC13a t 973°K. and one atmosphere. The e.m.f., denoted by E, was determined a t various compositions of the electrolytic solutions as a function of temperature. After having been corrected to that of one atmosphere of chlorine gas (corrections ranged from 0.010 to 0.012 v.) and for thermoelectric effects in the external circuit (corrections ranged from 0.005 to 0.008 v.), these data were represented (least squares fit) by linear equations of the form exl = ax,
+ bx,T
where axl and bx, are constants for the solution of PuCla having mole fraction X1 and T denotes the absolute temperature. The experimental points fit the linear equation with a standard deviation of 0.001 to 0.002 v. The results are given analytically in Table I.
1057
A&*. The relative partial molar free energy of mixing for supercooled liquid PuC4 a t 973°K. is given in column 5 of the same table. The standard molar quantities for the formation for solid PuCl3 at 973°K. are given in the last row of Table 111. These quantities were obtained by combining the reactions (a) and (b) in Table 111. The data for reaction (b), the freezing of liquid PuCla a t 973”K., were obtained from Table IV of reference 1. TABLE TI THE MOLARTHERMODYNAMIC QUANTITIES O F FORMATION OF LIQUID PuCls AND THE PARTIAL MOLAR FREEENERGY OF PuCls FOR THE SYSTEM PuCkNaCl at 973°K. XI
1 .ooo 0.750 .700
AFI, kcal.
AS,, e.u.
AHI, kcal.
FP. kcal.
- 169
-36.9 -36.5 -36.3 -36.1 -35.1
- 205 - 20G - 207 - 208 - 208
0 -2 -3 -4 -5
AS, e.u.
AH, kcal.
Ref.
-36.9 -14.7
-205 -15.3
This work 6
-51.6
-220
- 171 - 172 - 173
.600
.500 - 174 TABLE I TABLEI11 POTENTIOMETRIC DATAAS A FUNCTION OF TEMPERATURE THE STANDARD MOLAR THERMODYNAMIC QUANTITIES FOR FOR THE FORMATION OF PuCls IN THE LIQUID BINARY SYSTHE FORMATION OF PURE PuClt IN THE SUPERCOOLED TEM PuCls-KaC1 AT VARIOUS COMPOSITIONS LIQUID AND IN THE SOLID STATE No. . of ~. ~
21
c.
Max. dev., mv.
v.
0.750 2.984 - 0.5280 X T 1 2 .700 2.996- .5242 X 10-aT 1 1 .600 3.002 - .5220 X T 1 4 .500 3.004- ,5069 X 10-T f 3
Temp.
e.m.f.
te%n.
r%Xel
974 to 1028 22 945t0978 5 913 to 1020 16 945t01017 9
The calculations involve, first, extrapolation of the data to the state of pure PuC4 at 973°K. in order to obtain the thermodynamic quantities for the formation of the supercooled liquid salt. Using these results, the relative partial molar free energies of mixing liquid PuC4 were computed. Finally, using these data and data taken from the literature, the standard formation quantities for solid PuC4 were computed. Further details of these calculations are described below. The free energy (AFl), entropy (AS1) and enthalpy (AH1) of formation of PuC13 in the liquid PuC13-NaC1 solutions a t 973°K. for the various compositions were computed using the formulas AFt A& =
-
-3FexI
(4)
(& AFl)
(5)
AH1 = AFI
+ 973AS1
(6)
The results are shown in Table 11. The value of the free energy of formation of pure supercooled liquid PuC4, which will be denoted AF1*, represents the arithmetic average of the values obtained upon extrapolating the function FlE/XZ2 for the various solutions to the state X1 = 1.000. FIE denotes the excess partial free energy of mixing PuCls, Le., FIE = AF1 - AF1* - RTInX1. Theentropy of formation of pure supercooled liquid PuCh, A&*, was taken as the arithmetic average of the excess entropy of formation of PuC4, ASIE = AXl R In XI, for the solutions. The enthalpy of formation of pure supercooled liquid PuCL was obtained using the relation AHI* = AF1* 973.
+
+
Reaction at 973OK. 3 (a) Punia) 5 Clnw =
+
PUClrPi,*,
(b) PuClaniw (c)
PUCli,,
= PuClrc.
AFs kcal. -169 -1.0
+ 3 Clz(X, =
PuClac.)
-170
Reliability.-The use of an inert-porous-thoria partition for the purpose of reducing the rate of mutual dissolution of the metallic plutonium electrode and the electrolyte has been discussed.’ The necessary conditions for reversibility were satisfied by all the galvanic cells, Le., the e.m.f. was independent of time, independent of the direction of approach to temperature equilibrium and reproducible after passing a small electrodissolving and electrodepositing current through the cell. Electrolytic solutions of mole fraction X1 < 0.5 showed a weight gain after having been saturated with chlorine gas at 700” and quenched to room temperature. No data are reported for this range of composition since some plutonium must exist in an oxidation state higher than three and cells containing such electrolytes involve a mixed cell reaction. The C1:Pu ratio in solutions of mole fractions X1 > 0.5, after saturation with chlorine gas, was found to be 3.0 f 0.1 by chemical analysis. I t is concluded that the concentration of a higher oxidation state is negligible for the reported data. Discussion The standard free energy and entropy of formation of pure solid PuC4 have been determined to be, respectively, - 170 kcal./mole and -52 e.u. which agree with previous results‘ within the precision of the measurements. The PuC13-NaC1 system exhibits negative deviations from Roult’s law. (6) G. L. Brewer, L. Bromley, P. W. Gillea and N. L. Lofgren, “The Transuranic Elemente,” National Nuclear Energy Series, Div. IV, Vol. 14B, McGraw-Hill New York, N. Y., 1949, pp. 861-886.
NOTES
1058
Vol. 65
Acknowledgments.-We would like to thank R. D. Baker and W. J. Maraman for their interests
namely, 0.313 for fluorides, 0.318 for chlorides, 0.325 for bromides, and 0.332 for iodide^.^ The which have been an encouragement throughout the equations for the energy of formation of the alkaprogress of Ihe work. We are indebted to J. W. line earth halides a t their equilibrium internuclear :lnderson aiid A. N. Morgan for the machined distances from the infinitely separated ions are plutonium metal, to C. F. Metz, C. T . Apel, D. C. Croley and G. R. Waterbury for the chemical analyses and to R. E. Cowan and S. D. Stoddard 3.5e2 4 9 e b A - P- esp for the thoria crucibles. =o = +7 r0 + Thus a t the equilibrium internuclear distance ro ENERGIES OF THE GASEOCS ALKALIKE wg = - 3.je2 TO (1 - L?) (1 (1) EARTH HALIDES' Lysing the ionic polarizabilities given b y Tessmaii: BY DANIELCVBICCIOTTI Kahn and Shockley6 the energies calculated are binnford Rcsearch Instatute Menlo Park Cali fornia given in the first column of Table I.
(- y )
(g)ro
-lg
Recezued iyouember 1 B 1960
The binding energies of the gaseous alkali halides have been calculated successfully by Berkowitz2 with the method of Rittner.3 Since the internuclear distmces in the gaseous alkaline earth halides have been measured, it seemed useful to attempt to calculate their binding energies. In the method of Rittner the expression for the energy consists of terms for coulombic attraction polarization and "overlap" repulsion. For the alkaline earth halide molecules, which are assumed to be linear, the coulombic term consists of the attrai-tions of the positive ion for the two negative ions and the repulsion between the negative ions. The coulombic energy becomes ( - 3 . L 7 T ) , where r is the cation-anion distance. In each anion a dipole is induced by the field of the positive ion rind the other anion. The field at the anion center is (7e/4r2), and the dipole moment induced is (1 (its polarizability) times the field. The total energy contributed by this polarization is cy ( 7 t 1 4 r ~ times ) ~ 2 (because there are two anions) times l 2 (bemuse half of the energy is stored in the dipoles). I n the present treatment the dipoledipole energy and the van der Waals dispersion cnergy are neglected because they are small, amounting at most to one-half per cent. of the total energy The overlap repulsioii energy term is of ail empirical nature. The form used by Rittiier and Berkowitz for the alkali halides was A exp(-T/p). The parameter p was evaluated from the vibration frequeiicies of several gaseous alkali halides. It was found to be approximately the same for all the salts and roughly equal to the corresponding parameter for the crystalline solids. The parameter A vias evaluated for each salt by using the condition thitt the energy be at a minimum a t the equilibrium internuclear distance. For the alkaline earth halides there is not sufficient information on the vibration frequencies to evaluatr the parameter p . Therefore, the values determined for the solid alkali halides were used, (1) This research was supported by t h e United States Air Force through the Air Force Office of Scientific Research of the Air Research and Development Command undei Contract No AF 49(638)-89 Reproduction in nhole or in p a r t is permitted for any purpose of t h e United States Goberiinient ( 2 ) J Rerkoxilz .I Chsm Phys 29, 1386 (1958) (3) E S Rittner, z h z d , 19, lOJ0 (1951) 4hishin and V P Spiridonov Krrslallograha ( U S S R ) , 1, 475 (1958).
?)
TABLEI ENERGIES (KCAL./MOLE) OF THE REACTIONRIX*(g) = Rf++(g) 2X-(g) FOR THE ALKALINEEARTHHALIDES
+
BeFz Cl? B r2 12
MgFz Clz Br? I? CaF2 Cl? Br? 12
SrFl Clz Br2
I? BaF2 Cln Bn 12
Calcd. b y eq. 1
Expt.
Calcd. by eq. 2
662 623 606 565 560 511 490 473 484 442 426 409 465 416 402 386 445 393 377 363
748 877 662 64 1 589 555 530 519 502 475 458 453 482 450 431 420 436 428 407 389
748 70; 675 624 5')1 533 513 500 502 455 438 425 4x2 427 -113 400 463 -104 389 37i
The experimental values for the energy of formation from the infinitely separated ions are also given in Table I. They were taken from information from the Xational Bureau of Standards' plus the estimated heats of sublimation of t8healkaline earth halides of Brewer.* The agreement betn-een the experiment'al and calculat,ed values i in general is well outside the lyOerror o ment of the internuclear distance aiid the =t 10 to 20 kcal./mole which is the probable range of accuracy of the experimental heats. Because of this lack of agreement a different set of assumptions was tried for t8heoverlap repulsioii t,erm. The term was assumed to have the alternate form Ar-" in which t'he values of n given by Paulingg mere used. This set of n's was pre(5) D. Cubicciotti, J . C h e m . P h y s . , 31, l(i4ii (1959); 3 3 , 1579 (1960). H. Kahn a n d 13.. Shockley, Phus. Rev., 92, 890 (1953). (7) (a) F. D. Rossini, D. D. M'agman, 13.. H. Evans, S. Levine and I. Jaffe, National Bureau of Standards, Circular 500, 1952; (b) National Bureau of Standards Report 6484, 1959. (8) L. Brewer, "Chem. a n d Met. of RIisc. bIatls.," K N E S IV-19B McGraw-Hill Book Co., New York, N. T., 1950. (9) L. Pauling, "The Xature of the Chemical Bond." Cornel1 Press, Ithecrt, N. Y.,1960, 3rd Ed., Table 13-2, p. 509.
